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1 Charles Page High School “The Periodic Table”Adapted fromPre-AP ChemistryCharles Page High SchoolStephen L. Cotton
2 Section 6.1 Organizing the Elements OBJECTIVES:Explain how elements are organized in a periodic table.
3 Section 6.1 Organizing the Elements OBJECTIVES:Compare early and modern periodic tables.
4 Section 6.1 Organizing the Elements OBJECTIVES:Identify three broad classes of elements.
5 Section 6.1 Organizing the Elements A few elements, such as gold and copper, have been known for thousands of years - since ancient timesYet, only about 13 had been identified by the year 1700.As more were discovered, chemists realized they needed a way to organize the elements.
6 Section 6.1 Organizing the Elements Chemists used the properties of elements to sort them into groups.In 1829 J. W. Dobereiner arranged elements into triads – groups of three elements with similar propertiesOne element in each triad had properties intermediate of the other two elements
7 Mendeleev’s Periodic Table By the mid-1800s, about 70 elements were known to existDmitri Mendeleev – a Russian chemist and teacherArranged elements in order of increasing atomic massThus, the first “Periodic Table”
8 He left blanks for yet undiscovered elements MendeleevHe left blanks for yet undiscovered elementsWhen they were discovered, he had made good predictionsBut, there were problems:Such as Co and Ni; Ar and K; Te and I
9 A better arrangementIn 1913, Henry Moseley – British physicist, arranged elements according to increasing atomic numberThe arrangement used todayThe symbol, atomic number & mass are basic items included-textbook page 162 and 163
12 The Periodic Law says:When elements are arranged in order of increasing atomic number, there is a periodic repetition of their physical and chemical properties.Horizontal rows = periodsThere are 7 periodsVertical column = group (or family)Similar physical & chemical prop.Identified by number & letter (IA, IIA)
13 Areas of the periodic table Three classes of elements are: ) metals, 2) nonmetals, and ) metalloidsMetals: electrical conductors, have luster, ductile, malleableNonmetals: generally brittle and non-lustrous, poor conductors of heat and electricity
14 Areas of the periodic table Some nonmetals are gases (O, N, Cl); some are brittle solids (S); one is a fuming dark red liquid (Br)Notice the heavy, stair-step line?Metalloids: border the line-2 sidesProperties are intermediate between metals and nonmetals
15 Section 6.2 Classifying the Elements OBJECTIVES:Describe the information in a periodic table.
16 Section 6.2 Classifying the Elements OBJECTIVES:Classify elements based on electron configuration.
17 Section 6.2 Classifying the Elements OBJECTIVES:Distinguish representative elements and transition metals.
18 Squares in the Periodic Table The periodic table displays the symbols and names of the elements, along with information about the structure of their atoms:Atomic number and atomic massBlack symbol = solid; red = gas; blue = liquid (from the Periodic Table on our classroom wall)
20 Groups of elements - family names Group IA – alkali metalsForms a “base” (or alkali) when reacting with water (not just dissolved!)Group 2A – alkaline earth metalsAlso form bases with water; do not dissolve well, hence “earth metals”Group 7A – halogensMeans “salt-forming”
21 Electron Configurations in Groups Elements can be sorted into 4 different groupings based on their electron configurations:Noble gasesRepresentative elementsTransition metalsInner transition metalsLet’s now take a closer look at these.
22 Electron Configurations in Groups Noble gases are the elements in Group 8A (also called Group18 or 0)Previously called “inert gases” because they rarely take part in a reaction; very stable = don’t reactNoble gases have an electron configuration that has the outer s and p sublevels completely full
23 Electron Configurations in Groups Representative Elements are in Groups 1A through 7ADisplay wide range of properties, thus a good “representative”Some are metals, or nonmetals, or metalloids; some are solid, others are gases or liquidsTheir outer s and p electron configurations are NOT filled
24 Electron Configurations in Groups Transition metals are in the “B” columns of the periodic tableElectron configuration has the outer s sublevel full, and is now filling the “d” sublevelA “transition” between the metal area and the nonmetal areaExamples are gold, copper, silver
25 Electron Configurations in Groups Inner Transition Metals are located below the main body of the table, in two horizontal rowsElectron configuration has the outer s sublevel full, and is now filling the “f” sublevelFormerly called “rare-earth” elements, but this is not true because some are very abundant
26 Elements in the 1A-7A groups are called the representative elements outer s or p filling1A8A2A3A4A5A6A7A
27 The group B are called the transition elements These are called the inner transition elements, and they belong here
28 Group 1A are the alkali metals (but NOT H) Group 2A are the alkaline earth metalsH
29 Group 8A are the noble gases Group 7A is called the halogens
30 1s11s22s11s22s22p63s11s22s22p63s23p64s11s22s22p63s23p64s23d104p65s11s22s22p63s23p64s23d104p65s24d10 5p66s11s22s22p63s23p64s23d104p65s24d105p66s24f145d106p67s1H1Li3Na11K19Rb37Cs55Fr87Do you notice any similarity in these configurations of the alkali metals?
31 1s21s22s22p61s22s22p63s23p61s22s22p63s23p64s23d104p61s22s22p63s23p64s23d104p65s24d105p61s22s22p63s23p64s23d104p65s24d10 5p66s24f145d106p6HeDo you notice any similarity in the configurations of the noble gases?2Ne10Ar18Kr36Xe54Rn86
32 Elements in the s - blocks Alkali metals all end in s1Alkaline earth metals all end in s2really should include He, but it fits better in a different spot, since He has the properties of the noble gases, and has a full outer level of electrons.
33 Transition Metals - d block Note the change in configuration.s1 d5s1 d10d1d2d3d5d6d7d8d10
35 F - block f1 f5 f2 f3 f4 f6 f7 f8 f9 f10 f11 f12 f14 f13 Called the “inner transition elements”f1f5f2f3f4f6f7f8f9f10f11f12f14f13
36 Each row (or period) is the energy level for s and p orbitals. 1234567Period NumberEach row (or period) is the energy level for s and p orbitals.
37 The “d” orbitals fill up in levels 1 less than the period number, so the first d is 3d even though it’s in row 4.12345673d4d5d
38 f orbitals start filling at 4f, and are 2 less than the period number 12345674f5ff orbitals start filling at 4f, and are 2 less than the period number
39 Section 6.3 Periodic Trends OBJECTIVES:Describe trends among the elements for atomic size.
40 Section 6.3 Periodic Trends OBJECTIVES:Explain how ions form.
41 Section 6.3 Periodic Trends OBJECTIVES:Describe periodic trends for first ionization energy, ionic size, and electronegativity.
42 Trends in Atomic SizeFirst problem: Where do you start measuring from?The electron cloud doesn’t have a definite edge.They get around this by measuring more than 1 atom at a time.
43 Atomic Size}RadiusMeasure the Atomic Radius - this is half the distance between the two nuclei of a diatomic molecule.
44 ALL Periodic Table Trends Influenced by three factors:1. Energy LevelHigher energy levels are further away from the nucleus.2. Charge on nucleus (# protons)More charge pulls electrons in closer. (+ and – attract each other)3. Shielding effect(blocking effect?)
45 What do they influence?Energy levels and Shielding have an effect on the GROUP ( )Nuclear charge has an effect on a PERIOD ( )
46 #1. Atomic Size - Group trends HAs we increase the atomic number (or go down a group). . .each atom has another energy level,so the atoms get bigger.LiNaKRb
47 #1. Atomic Size - Period Trends Going from left to right across a period, the size gets smaller.Electrons are in the same energy level.But, there is more nuclear charge.Outermost electrons are pulled closer.NaMgAlSiPSClAr
48 RbKPeriod 2NaLiAtomic Radius (pm)KrArNeH310Atomic Number
49 Ions Some compounds are composed of particles called “ions” An ion is an atom (or group of atoms) that has a positive or negative chargeAtoms are neutral because the number of protons equals electronsPositive and negative ions are formed when electrons are transferred (lost or gained) between atoms
50 Metals tend to LOSE electrons, from their outer energy level IonsMetals tend to LOSE electrons, from their outer energy levelSodium loses one: there are now more protons (11) than electrons (10), and thus a positively charged particle is formed = “cation”The charge is written as a number followed by a plus sign: Na1+Now named a “sodium ion”
51 Nonmetals tend to GAIN one or more electrons IonsNonmetals tend to GAIN one or more electronsChlorine will gain one electronProtons (17) no longer equals the electrons (18), so a charge of -1Cl1- is re-named a “chloride ion”Negative ions are called “anions”
52 #2. Trends in Ionization Energy Ionization energy is the amount of energy required to completely remove an electron (from a gaseous atom).Removing one electron makes a 1+ ion.The energy required to remove only the first electron is called the first ionization energy.
53 Ionization EnergyThe second ionization energy is the energy required to remove the second electron.Always greater than first IE.The third IE is the energy required to remove a third electron.Greater than 1st or 2nd IE.
54 Symbol First Second Third Table 6.1, p. 173Symbol First Second ThirdHHeLiBeBCNO F Ne
55 Symbol First Second Third HHeLiBeBCNO F NeWhy did these values increase so much?
56 What factors determine IE The greater the nuclear charge, the greater IE.Greater distance from nucleus decreases IEFilled and half-filled orbitals have lower energy, so achieving them is easier, lower IE.Shielding effect
57 ShieldingThe electron on the outermost energy level has to look through all the other energy levels to see the nucleus.Second electron has same shielding, if it is in the same period
58 Ionization Energy - Group trends As you go down a group, the first IE decreases because...The electron is further away from the attraction of the nucleus, andThere is more shielding.
59 Ionization Energy - Period trends All the atoms in the same period have the same energy level.Same shielding.But, increasing nuclear chargeSo IE generally increases from left to right.Exceptions at full and 1/2 full orbitals.
60 He has a greater IE than H. Both elements have the same shielding since electrons are only in the first levelBut He has a greater nuclear chargeHFirst Ionization energyAtomic number
61 These outweigh the greater nuclear charge Li has lower IE than Hmore shieldingfurther awayThese outweigh the greater nuclear chargeHFirst Ionization energyLiAtomic number
62 greater nuclear charge HeBe has higher IE than Lisame shieldinggreater nuclear chargeFirst Ionization energyHBeLiAtomic number
63 greater nuclear charge HeB has lower IE than Besame shieldinggreater nuclear chargeBy removing an electron we make s orbital half-filledFirst Ionization energyHBeBLiAtomic number
64 First Ionization energy HeFirst Ionization energyHCBeBLiAtomic number
65 First Ionization energy HeNFirst Ionization energyHCBeBLiAtomic number
66 HeOxygen breaks the pattern, because removing an electron leaves it with a 1/2 filled p orbitalNFirst Ionization energyHCOBeBLiAtomic number
67 First Ionization energy HeFNFirst Ionization energyHCOBeBLiAtomic number
68 Ne has a lower IE than He Both are full, Ne has more shielding Greater distanceFNFirst Ionization energyHCOBeBLiAtomic number
69 Na has a lower IE than Li Both are s1 Na has more shielding HeNeNa has a lower IE than LiBoth are s1Na has more shieldingGreater distanceFNFirst Ionization energyHCOBeBLiNaAtomic number
71 Driving ForcesFull Energy Levels require lots of energy to remove their electrons.Noble Gases have full orbitals.Atoms behave in ways to try and achieve a noble gas configuration.
72 2nd Ionization EnergyFor elements that reach a filled or half-filled orbital by removing 2 electrons, 2nd IE is lower than expected.True for s2Alkaline earth metals form 2+ ions.
73 3rd IEUsing the same logic s2p1 atoms have an low 3rd IE.Atoms in the aluminum family form 3+ ions.2nd IE and 3rd IE are always higher than 1st IE!!!
74 Trends in Ionic Size: Cations Cations form by losing electrons.Cations are smaller than the atom they came from – not only do they lose electrons, they lose an entire energy level.Metals form cations.Cations of representative elements have the noble gas configuration before them.
75 Ionic size: Anions Anions form by gaining electrons. Anions are bigger than the atom they came from – have the same energy level, but a greater area the nuclear charge needs to coverNonmetals form anions.Anions of representative elements have the noble gas configuration after them.
76 Configuration of IonsIons always have noble gas configurations ( = a full outer level)Na atom is: 1s22s22p63s1Forms a 1+ sodium ion: 1s22s22p6Same configuration as neon.Metals form ions with the configuration of the noble gas before them - they lose electrons.
77 Configuration of IonsNon-metals form ions by gaining electrons to achieve noble gas configuration.They end up with the configuration of the noble gas after them.
78 Ion Group trends Each step down a group is adding an energy level Ions therefore get bigger as you go down, because of the additional energy level.Li1+Na1+K1+Rb1+Cs1+
79 Ion Period TrendsAcross the period from left to right, the nuclear charge increases - so they get smaller.Notice the energy level changes between anions and cations.N3-O2-F1-B3+Li1+Be2+C4+
80 Size of Isoelectronic ions Iso- means “the same”Isoelectronic ions have the same # of electronsAl3+ Mg2+ Na1+ Ne F1- O2- and N3-all have 10 electronsall have the same configuration: 1s22s22p6 (which is the noble gas: neon)
81 Size of Isoelectronic ions? Positive ions that have more protons would be smaller (more protons would pull the same # of electrons in closer)N3-O2-F1-NeNa1+Al3+10987131211Mg2+
82 #3. Trends in Electronegativity Electronegativity is the tendency for an atom to attract electrons to itself when it is chemically combined with another element.They share the electron, but how equally do they share it?An element with a big electronegativity means it pulls the electron towards itself strongly!
83 Electronegativity Group Trend The further down a group, the farther the electron is away from the nucleus, plus the more electrons an atom has.Thus, more willing to share.Low electronegativity.
84 Electronegativity Period Trend Metals are at the left of the table.They let their electrons go easilyThus, low electronegativityAt the right end are the nonmetals.They want more electrons.Try to take them away from othersHigh electronegativity.
85 The arrows indicate the trend: Ionization energy and Electronegativity INCREASE in these directions
86 Atomic size and Ionic size increase in these directions:
87 Summary Chart of the trends: Figure 6.22, p.178 End of Chapter 6